chemistry-class-11-unit-08-chapter-03-redox-reactions-l-3-3-uroz1kg7t20

### <Success style="font-size:25px"> Redox Reactions L-3 </Success>
### <primary style="font-size:24px"> Disproportionation reaction </primary>
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* Disproportionation is a definite type of redox reaction in which a specie is simultaneously reduced and oxidized to form two different products

* The UV photolysis of mercury(I) chloride

- $\mathrm{Hg}_2 \mathrm{Cl}_2 \rightarrow \mathrm{Hg}+\mathrm{HgCl}_2$

* It is a disproportionation reaction

* Chlorine can assume oxidation states of $-1,+1,+3,+5$, or +7 within the corresponding anions $\mathrm{Cl}^{-}, \mathrm{ClO}^{-}, \mathrm{ClO}_2^{-}$, $\mathrm{ClO}_3^{-}$, or $\mathrm{ClO}_4^{-}$, known commonly and respectively as chloride, hypochlorite, chlorite, chlorate, and perchlorate

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<footer style = "font-size: 18px">Redox reaction $\rightarrow$ Topics to be covered $\rightarrow$ <span style = "color:blue"> Disproportionation reaction </span> $\rightarrow$ Disproportionation reaction $\rightarrow$ Disproportionation reaction </footer>

### <Success style="font-size:25px"> Redox Reactions L-3 </Success>
### <primary style="font-size:24px"> Disproportionation reaction </primary>
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* Metal salts like $Hg_2Cl_2$ (Hg +1 oxidation state) Cu(I)Cl can goes to $Cu^0$ and $Cu^{2+}$ this is a typical dipproponation reaction

* $Cu^{+1}$ (aq) ${\xrightarrow{CH_3CN}}$ $Cu(CH_3CN)_4^+$ $\rightarrow$ Solid state ion aqalous

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<footer style = "font-size: 18px">Topics to be covered $\rightarrow$ Disproportionation reaction $\rightarrow$ <span style = "color:blue"> Disproportionation reaction </span> $\rightarrow$ Disproportionation reaction $\rightarrow$ Disproportionation reaction </footer>

### <Success style="font-size:25px"> Redox Reactions L-3 </Success>
### <primary style="font-size:24px"> Disproportionation reaction </primary>
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- $\begin{aligned} & \mathrm{P}_4^{\circ}+3 \mathrm{HO}^{-}+3 \mathrm{H}_2 \mathrm{O} \longrightarrow \stackrel{}{\mathrm{PH}_3^{-3}}+2 \mathrm{H}_2^{+1} \mathrm{PO}_2^{-} \\ & \mathrm{S}_8^0+12 \mathrm{HO}^{-} \longrightarrow 4^{-2} \mathrm{~S}^{2-}+2 \mathrm{~S}_2^2 \mathrm{O}_3^{2-}+6 \mathrm{H}_2 \mathrm{O} \\ & \mathrm{Cl}_2^0+2 \mathrm{HO}^{-} \longrightarrow \mathrm{Cl}^{-}+\mathrm{ClO}^{-}+\mathrm{H}_2 \mathrm{O} \\ &  \end{aligned}$

- $2F_2^o(g)+2HO^-(aq.) \rightarrow 2F^-(aq.)+OF_2(g)+H_2O (l)$

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<footer style = "font-size: 18px">Disproportionation reaction $\rightarrow$ Disproportionation reaction $\rightarrow$ <span style = "color:blue"> Disproportionation reaction </span> $\rightarrow$ Disproportionation reaction $\rightarrow$ Redox titrations </footer>

### <Success style="font-size:25px"> Redox Reactions L-3 </Success>
### <primary style="font-size:24px"> Disproportionation reaction </primary>
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* Chlorine gas reacts with dilute sodium hydroxide to form sodium chloride, sodium chlorate and water

- $3 \mathrm{Cl}_2+6 \mathrm{OH}^{-} \rightarrow 5 \mathrm{Cl}^{-}+\mathrm{ClO}_3^{-}+3 \mathrm{H}_2 \mathrm{O}$

* Bromine fluoride undergoes disproportionation reaction to bromine trifluoride and bromine

- $3BrF \rightarrow BrF_3+Br_2$

* Nitrogen has oxidation state +IV in nitrogen dioxide, but when this compound reacts with water, it forms nitric acid and nitrous acid

- $2 \mathrm{NO}_2+\mathrm{H}_2 \mathrm{O} \rightarrow \mathrm{HNO}_3+\mathrm{HNO}_2$

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<footer style = "font-size: 18px">Disproportionation reaction $\rightarrow$ Disproportionation reaction $\rightarrow$ <span style = "color:blue"> Disproportionation reaction </span> $\rightarrow$ Redox titrations $\rightarrow$ Balancing the reaction </footer>

### <Success style="font-size:25px"> Redox Reactions L-3 </Success>
### <primary style="font-size:24px"> Redox titrations </primary>
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* Redox reactions are very useful in quantitative analysis by redox titrations These titrations involve the reactions between oxidizing and reducing agents and help in estimating the amount of unknown substances present in the solutions

* Reagents-oxidizing and reducing

* Potassium permanganate $\mathrm{KMnO}_4$ : Permanganometry ceric sulfate / Ceric ammonium sulfate $\mathrm{Ce}\left(\mathrm{SO}_4\right)_2 \cdot 2\left(\mathrm{NH}_4\right)_2 \mathrm{SO}_4 \cdot 4 \mathrm{H}_2 \mathrm{O}$ : Crimetry potassium dichromate $\mathrm{K}_2 \mathrm{Cr}_2 \mathrm{O}_7$ : Dichromometry iodine $\mathrm{I}_2$ : iodimetry, iodometry

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<footer style = "font-size: 18px">Disproportionation reaction $\rightarrow$ Disproportionation reaction $\rightarrow$ <span style = "color:blue"> Redox titrations </span> $\rightarrow$ Balancing the reaction $\rightarrow$ Redox titrations </footer>

### <Success style="font-size:25px"> Redox Reactions L-3 </Success>
### <primary style="font-size:24px"> Balancing the reaction </primary>
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* Assign A, B, C, D
* Assignment of oxidation state
* Calculation of increase and decrease in oxidation number per element
* Write the overall reaction

* $SO_2 (reducing) +H_2O \rightarrow H_2SO_3 \rightarrow SO_3^{2-} \rightarrow SO_4^{2-}$

* $Cr_2O_7^{2-}+3SO_3^{2-} \rightarrow 2Cr^{3+}+3SO_4^{2-}$

* $HNO_3+3H_2S \rightarrow 2NO+S^{0}$

* $KMnO_4+KCl \rightarrow Mn^{+2}SO_4+Cl_2$

* $2KMnO_4+10KCl+8H_2SO_4 \rightarrow 2MnSO_4+5Cl_2+8H_2O+6K_2SO_4$

* $K_3Fe(CN)_6+Cr_2^{+3}O_3 /K_4Fe(CN)_6+CrO_4^{2-}$

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<footer style = "font-size: 18px">Disproportionation reaction $\rightarrow$ Redox titrations $\rightarrow$ <span style = "color:blue"> Balancing the reaction </span> $\rightarrow$ Redox titrations $\rightarrow$ Applications of permanganometry </footer>

### <Success style="font-size:25px"> Redox Reactions L-3 </Success>
### <primary style="font-size:24px"> Redox titrations </primary>
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* $\mathrm{KMnO}_4$ Powerful oxidant that the most widely used

* In strongly acidic solutions ( $1 \mathrm{M} \quad \mathrm{H}_2 \mathrm{SO}_4$ or $\mathrm{HCl}, \mathrm{pH}<1$ )

* $\mathrm{MnO}_4^{-}(violet)+8 \mathrm{H}^{+}+5 \mathrm{e}^{-}=\mathrm{Mn}^{2+}(colorless)+4 \mathrm{H}_2 \mathrm{O}$

* $\mathrm{E}^0=1.51 \mathrm{~V}$

* $\mathrm{KMnO}_4$ is a self-indicator

* Potassium permanganate is not primary standard, because traces of $\mathrm{MnO}_2$ are invariably present

* Standardization by titration of sodium oxalate (primary standard)-
- $\begin{array}{r}
2 \mathrm{KMnO}_4+5 \mathrm{Na}_2(\mathrm{COO})_2+8 \mathrm{H}_2 \mathrm{SO}_4=2 \mathrm{MnSO}_4+\mathrm{K}_2 \mathrm{SO}_4+ \\
5 \mathrm{Na}_2 \mathrm{SO}_4+10 \mathrm{CO}_2+8 \mathrm{H}_2 \mathrm{O}
\end{array}$

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<footer style = "font-size: 18px">Redox titrations $\rightarrow$ Balancing the reaction $\rightarrow$ <span style = "color:blue"> Redox titrations </span> $\rightarrow$ Applications of permanganometry $\rightarrow$ Redox titration </footer>

### <Success style="font-size:25px"> Redox Reactions L-3 </Success>
### <primary style="font-size:24px"> Applications of permanganometry </primary>
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* Determination of $\mathrm{H}_2 \mathrm{O}_2$

- $2 \mathrm{KMnO}_4+5 \mathrm{H}_2 \mathrm{O}_2+3 \mathrm{H}_2 \mathrm{SO}_4=2 \mathrm{MnSO}_4+\mathrm{K}_2 \mathrm{SO}_4+5 \mathrm{O}_2+8 \mathrm{H}_2 \mathrm{O}$

* Determination of $\mathrm{NaNO}_2$

- $2 \mathrm{NaNO}_2+\mathrm{H}_2 \mathrm{SO}_4=\mathrm{Na}_2 \mathrm{SO}_4+\mathrm{HNO}_2$

- $2 \mathrm{KMnO}_4+5 \mathrm{HNO}_2+3 \mathrm{H}_2 \mathrm{SO}_4=2 \mathrm{MnSO}_4+\mathrm{K}_2 \mathrm{SO}_4+5 \mathrm{HNO}_3 + 3 \mathrm{H}_2 \mathrm{O}$

* Determination of  $\mathrm{FeSO}_4 $

- $\begin{aligned}
&  2 \mathrm{KMnO}_4+10 \mathrm{FeSO}_4+8 \mathrm{H}_2 \mathrm{SO}_4=2 \mathrm{MnSO}_4+\mathrm{K}_2 \mathrm{SO}_4+ \\
& 5 \mathrm{Fe}_2\left(\mathrm{SO}_4\right)_3+8 \mathrm{H}_2 \mathrm{O} \\
&
\end{aligned}$

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<footer style = "font-size: 18px">Balancing the reaction $\rightarrow$ Redox titrations $\rightarrow$ <span style = "color:blue"> Applications of permanganometry </span> $\rightarrow$ Redox titration $\rightarrow$ Redox titation </footer>

### <Success style="font-size:25px"> Redox Reactions L-3 </Success>
### <primary style="font-size:24px"> Redox titration </primary>
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* Manganese(II) reacts with sodium bismuthate

- $\left.\mathrm{Mn}^{2+} (\mathrm{aq}\right)+\mathrm{NaBiO}_3(\mathrm{~s}) \rightarrow \mathrm{Bi}^{3+} \mathrm{aq}+\mathrm{MnO}_4^{-}(\mathrm{aq})$

* Oxidation

- $4 \mathrm{H}_2 \mathrm{O}(1)+\mathrm{Mn}^{2+}(\mathrm{aq}) \rightarrow \mathrm{MnO}_4$ (aq) $+8 \mathrm{H}^{+}(\mathrm{aq})+5 \mathrm{e}$

* Reduction

- $2 \mathrm{e}^{-}+6 \mathrm{H}^{+}+\mathrm{BiO}_3^{-} \rightarrow \mathrm{Bi}^{3+}(\mathrm{aq})+3 \mathrm{H}_2 \mathrm{O}(1)$

- $14 \mathrm{H}^{+}(\mathrm{aq})+2 \mathrm{Mn}^{2+}(\mathrm{aq})+5 \mathrm{NaBiO}_3(\mathrm{~s}) \rightarrow$
- $7 \mathrm{H}_2 \mathrm{O}(\mathrm{l})+2 \mathrm{MnO}_4^{-}(\mathrm{aq})+5 \mathrm{Bi}^{3+}(\mathrm{aq})+5 \mathrm{Na}^{+}(\mathrm{aq})$

* Thus in acidic media, $\mathrm{H}^{+}$ions and water are added to half-reactions to balance the overall reaction

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<footer style = "font-size: 18px">Redox titrations $\rightarrow$ Applications of permanganometry $\rightarrow$ <span style = "color:blue"> Redox titration </span> $\rightarrow$ Redox titation $\rightarrow$ Thank you </footer>

### <Success style="font-size:25px"> Redox Reactions L-3 </Success>
### <primary style="font-size:24px"> Redox titation </primary>
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* The reaction between potassium permanganate and sodium sulfite

- $2 \mathrm{KMnO}_4+3 \mathrm{Na}_2 \mathrm{SO}_3+\mathrm{H}_2 \mathrm{O} \rightarrow 2 \mathrm{MnO}_2+3 \mathrm{Na}_2 \mathrm{SO}_4+2 \mathrm{KOH}$

* In basic media, $\mathrm{HO}^{-}$ions and water are added to half reactions to balance the overall reaction

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<footer style = "font-size: 18px">Applications of permanganometry $\rightarrow$ Redox titration $\rightarrow$ <span style = "color:blue"> Redox titation </span> $\rightarrow$ Thank you $\rightarrow$  </footer>